Experiment 5 Discussion

Daria Evans Chemistry 1AL, TH 11/6/08 Experiment 5 The Molar Volume of Gases Discussion Part 1: Production of Oxygen Gas The purpose of part 1 in the experiment was to determine the volume of oxygen produced in a reaction through water displacement. For the procedure in part 1 of the experiment I connected a 600-mL beaker with water to a 500-mL Erlenmeyer flask filled with water with a rubber tube. On the other side I connected a 250-mL Erlenmeyer flask with 20ml of 3% H2O2 solution and a test tube of 4.5 mL of 3M FeCl3 inside to the large 500-mL Erlenmeyer flask. After equalizing the pressures inside and outside the beaker and Erlenmeyer flasks, I let the FeCl3 spill inside the flask and watched how the reaction occurred. The FeCl3 (catalyst) sped up the reaction and soon I felt heat being released from the reaction. As a result of the reaction the water in the beaker began to rise and I was able to measure the change in volume of the water and as a result the change in volume of oxygen gas. First I measured the mass of flask A and its contents before the reaction and got 143.89g and for the mass of flask A after the reaction I got 143.582g. Afterwards, I used a thermometer and measured the temperature of flask A (24.2C) and the temperature of flask B (24.8C). I calculated the average of the temperatures to be 24.5C. By using the formula PV=nRT I was able to figure out the volume of oxygen collected (220.5mL) and combine the standard temperature and pressure to the amounts I calculated in the experiment. The barometric pressure during the time was 1.01 atm or 767.6 torr. Also, by using Appendix 1 I figured out the water vapor pressure at the average temperature of the gas to be 23.0665 torr. From the reaction I came up witht the following balanced equation: 2H2O2 (l)

2H2O (l) + O2(g)

I was able to calculate the moles of oxygen (0.009625 mol) and the average temperature of gas (297.5K). From the rest of my data I was able to determine the pressure of oxygen (744.5 torr), volume of oxygen (20.6 L) and molar volume of O2 (22.91 L/mol). This value is relatively close to the standard volume of 22.4 L/mol. Therefore, my data is accurate. Part 2: The Molar Volume of Hydrogen The purpose of part 2 of the experiment was to determine the volume of hydrogen gas produced in the reaction through water displacement. I filled up my 600-mL beaker with water and connected it to the 500-mL Erlenmeyer flask and then connected another 250-mL Erlenmeyer flask with a 0.066g magnesium metal sample with a small test tube of 4.5 mL of 3M HCl solution to the

Daria Evans Chemistry 1AL, TH 11/6/08 same 500-mL flask. Then, I calculated the temperature of the gas in flask A (23.5C) and the temperature in flask B (26.9C). Their average temperatures were 25.2C. I measured the volume of hydrogen gas collected to be 39.8 mL and the barometric pressure was 1.01 atm. By using Appendix 2 I calculated the water vapor pressure at the average temperature of the gas to be 24.0466 torr. From my data and the balanced chemical equation of the reaction I determined the moles of hydrogen (.0027 mol). Mg (s) + 2 HCl (aq)

MgCl2 (aq) + H2 (g)

The average temperature of gas was 298.2 K and the pressure of hydrogen was calculated to be 743.55 torr. By using the formula PV=nRT I was able to calculate the volume of H2 at STP to be 13.203 L and the molar volume I calculated to be 14.74 L/mol. These values are smaller than 22.4 L/mol that is expected because of a particular source of error. Our source or error consisted of the fact that not enough water was displaced since the stopper on the 250-mL Erlenmeyer flask was not tightly closed off and therefore not enough water actually pushed through the rubber tubes to get an accurate reading. We discovered that by using water displacement we were able to measure the amount of oxygen or hydrogen gas we produced in both reactions. And by using stoichiometry we were able to test our results and compare them to STP.